Chapter 1 Notes – Matter & Measurements

1.1 Chemistry: The Central Science

Chemistry is called “the central science” because it underpins virtually every other scientific discipline (biology, physics, geology, medicine, environmental science, etc.).
Chemistry = study of the nature, properties, and transformations of matter.
Matter = anything that has mass and occupies space; tangible (can be seen, touched, tasted, smelled).
Scientific Method = cyclical process of observation → hypothesis → experimentation → theory refinement; basis for all chemical knowledge.
Property (used for identification):
- Size, color, temperature, chemical composition, chemical reactivity, etc.
Two kinds of change:
- Physical change: no change in chemical identity (e.g., melting ice, sugar dissolving in water).
- Chemical change: alters chemical makeup; new substance(s) with different properties (e.g., sugar → caramel on heating).
- Worked Example 1.1:
  • (a) Sugar + water → physical (can evaporate water, recover sugar).
  • (b) Heating sugar → caramel → chemical (new color, consistency).
Table 1.1 highlights contrasting physical/chemical properties of water, sucrose, and sodium bicarbonate (e.g., ext{mp}{\text{H}2\text{O}} = 0\,^{\circ}!\text{C} vs. sucrose decomposition at 160\,^{\circ}!\text{C}).
Visual concept map shows chemistry’s bridges to biochemistry, physics, medicine, agriculture, etc.

1.2 States of Matter

Three physical states:
• Solid (s): definite shape & volume.
• Liquid (l): definite volume, takes shape of container.
• Gas (g): neither definite volume nor shape.
Change of state = conversion between s, l, g (melting, boiling, condensation, sublimation, deposition, freezing).
Example (Worked Ex. 1.2): Formaldehyde mp = -92\,^{\circ}!\text{C}, bp = -19.5\,^{\circ}!\text{C}, so at room T (25 °C) it is a gas (25 °C > bp).

1.3 Classification of Matter

Matter → either a pure substance or a mixture.
Pure substance: uniform composition; constant properties.
• Element: cannot be chemically broken down further.
• Compound: can be decomposed to simpler substances by chemical reaction.
Mixture: physical blend of ≥ 2 substances, each retaining identity.
• Homogeneous (solution): uniform composition (air, salt water, coffee).
• Heterogeneous: non-uniform (pot pie, concrete, chocolate-chip cookies).
Concept Map and decision tree illustrate classification path (pure vs mixture, homogeneous vs heterogeneous, element vs compound).
Worked Example 1.3:
• Vanilla ice cream → homogeneous mixture.
• Sugar → pure substance, compound (made of C, H, O atoms).

1.4 Chemical Elements and Symbols

118 known elements (≈91 natural).
Symbol rules: first letter uppercase, second lowercase (Na, Fe, Cl). Some derive from Latin names (Na = natrium, Fe = ferrum, Au = aurum).
Table 1.2 lists common element symbols (Al, He, Cu, etc.).
Abundance: Earth’s crust: O 46.1 %, Si 28.2 %, Al 8.2 % …; Human body: O 61 %, C 23 %, H 10 %, N 2.6 %, etc.
Chemical formula = symbol notation with subscripts denoting number of atoms (e.g., \text{H}_2\text{O}). No subscript → understood 1.
Table 1.4: Elements essential to life; macro-nutrients (> 100 mg/day, e.g., Ca, P, K, Na, Mg, S) vs micro-nutrients (≤ 15 mg/day, e.g., Fe, Zn, I).

1.5 Chemical Reactions (Examples of Chemical Change)

Chemical reaction = process changing identity/composition of substance(s).
• Reactants (left) → Products (right); arrow indicates change & conditions.
Water electrolysis: 2\,\text{H}2\text{O}(l) \xrightarrow[\text{electric current}]{} 2\,\text{H}2(g) + \text{O}_2(g) (gas bubbles).
Nickel + HCl(aq): color change + gas → \text{Ni}(s)+2\,\text{HCl}(aq)→\text{NiCl}2(aq)+\text{H}2(g).

1.6 Physical Quantities, Units & Scientific Notation

Physical quantity = measurable property described by number + unit.
SI base units: mass kg, length m, volume m^3, temperature K, time s.
Metric choices: gram (g) for mass, liter (L) for volume, Celsius (°C) for T.
Derived units: speed m/s; density g/cm^3.
Prefixes (Table 1.6): 10^6 = mega (M), 10^3 = kilo (k), 10^{-3} = milli (m), 10^{-6} = micro (\mu), 10^{-9} = nano (n), …
Scientific notation rules:
• Standard: N \times 10^{n}, where 1\le N
Worked Ex. 1.4: 0.000000120 m = 1.20 \times 10^{-7}\,\text{m} = 120 nm = 0.120 µm.

1.7 Measuring Mass, Length & Volume

Mass = amount of matter; Weight = gravitational force on that mass.
Typical conversions (Tables 1.7–1.9):
• 1\,\text{kg}=2.205\,\text{lb}; 1\,\text{lb}=454\,\text{g}.
• 1\,\text{m}=39.37\,\text{in}=1.0936\,\text{yd}.
• 1\,\text{L}=1000\,\text{mL}=1.057\,\text{qt}.

1.8 Measurement & Significant Figures

Any measurement has uncertainty; record all certain digits + one estimate.
Significant-figure (sig-fig) rules:
1. Interior zeros significant (94.072 g → 5 sig-figs).
2. Leading zeros NOT significant (0.0076 mL → 2 sig-figs).
3. Trailing zeros after decimal significant (138.200 m → 6 sig-figs).
4. Trailing zeros before implied decimal ambiguous (23 000 kg → 2–5 sig-figs; clarify with decimal or scientific notation).
Exact numbers (definitions, counted objects) have infinite sig-figs (e.g., 1 ft = 12 in exactly).
Scientific notation clarifies sig-figs (e.g., 2.30\times10^4 has 3 sig-figs).

1.9 Rounding Off Numbers

Multiplication/Division → answer limited to smallest # sig-figs among inputs.
Addition/Subtraction → answer limited to least digits right of decimal among inputs.
Rounding rules:
• First dropped digit ≤ 4 → truncate.
• First dropped digit ≥ 5 → round up.
Example 1.9: 124 lb + 1.884 lb = 125.884 lb → round to 126 lb (no decimals, because 124 lb had none).
Example 1.10: \frac{13.75\,\text{cups}}{18\,\text{cups}} = 0.763888… → 0.76 (2 sig-figs from 18).

1.10 Problem Solving – Unit Conversions

Factor-label (dimensional-analysis) method:
1. List given quantity + units.
2. Identify desired quantity + units.
3. Choose conversion factor(s) so unwanted units cancel.
4. Calculate, then check with ball-park estimate.
Conversion factor = ratio of equivalent quantities; numerically equal to 1, e.g., \frac{1000\,\text{mL}}{1\,\text{L}}.
Worked Exs.:
• 0.75 lb → g: 0.75\,\text{lb}\times\frac{454\,\text{g}}{1\,\text{lb}} = 3.4\times10^{2}\,\text{g}.
• 0.50 qt → dL: chain 0.50\,\text{qt}\times\frac{946\,\text{mL}}{1\,\text{qt}}\times\frac{1\,\text{dL}}{100\,\text{mL}}=4.7\,\text{dL}.
• Dosage problems (painkiller, digitalis) use concentration or body-weight factors.

1.11 Temperature, Heat & Energy

Energy (J) = capacity to do work/produce heat; Heat = energy transfer due to T difference; Temperature = measure of average heat energy.
Temperature scales & conversions:
• T{\text{K}} = T{\,^{\circ}!\text{C}} + 273.15.
• T{\,^{\circ}!\text{C}} = (T{\,^{\circ}!\text{F}} – 32) \times \frac{1}{1.8}.
• 1 °C change = 1 K change = 1.8 °F change.
Example 1.16: 107 °F → \frac{(107−32)}{1.8}=41.7\,^{\circ}!\text{C}.
Heat units & conversions:
• 1\,\text{cal}=4.184\,\text{J}; 1\,\text{kcal}=1000\,\text{cal}=4.184\,\text{kJ}.
• Nutritionists’ “Calorie” (Cal) = kcal.
Specific heat (c): energy to raise 1 g of substance by 1 °C; unique per material (Table 1.10).
Heat equation: q = m\,c\,\Delta T.
Worked 1.17 (bath): 95 kg water, \Delta T = 25\,^{\circ}!\text{C}.
q = (95\times10^{3}\,\text{g})(1.00\,\text{cal/g °C})(25 °C)=2.4\times10^{6}\,\text{cal}=1.0\times10^{7}\,\text{J}.

1.12 Density & Specific Gravity

Density: \rho = \frac{m}{V}, commonly g/cm^3 (solids) or g/mL (liquids). 1 mL = 1 cm^3.
Specific gravity (sg): \text{sg} = \frac{\text{density of substance}}{\text{density of water (at same T)}}; since \rho_{\text{water}} \approx 1\,\text{g/mL} at RT, sg ≈ numeric density.
Density trends: less dense materials float on denser fluids (ice floats on water; cork floats; gold sinks).
Hydrometer/urinometer measures sg of solutions; medical use: urine solids content.
Worked 1.18: find V for 25.0 g isopropanol (d = 0.7855 g/mL):
V = m\times\frac{1}{\rho}=25.0\,\text{g}\times\frac{1\,\text{mL}}{0.7855\,\text{g}} = 31.8\,\text{mL} (≈30 mL as estimated).

Conceptual Connections & Applications

Chemical & physical property distinctions guide material identification, safety, and reaction monitoring.
Accurate measurement (sig-figs, rounding) underlies quantitative chemistry (stoichiometry, solution prep, dosage calculation).
Temperature & heat principles link to thermodynamics, reaction energetics, physiology (fever thresholds, caloric content).
Density/specific gravity applicable to flotation, separation techniques, clinical diagnostics.
Unit conversion fluency ensures cross-disciplinary communication (engineering, medicine, environmental science) and prevents critical errors (e.g., medication dosing).

Key Formulas (collective)

q = m c \Delta T (heat change).
\rho = \frac{m}{V}; m = \rho V; V = \frac{m}{\rho}.
T{\text{K}} = T{\,^{\circ}!\text{C}} + 273.15; T{\,^{\circ}!\text{C}} = T{\text{K}} – 273.15.
T{\,^{\circ}!\text{F}} = (1.8)(T{\,^{\circ}!\text{C}}) + 32; T{\,^{\circ}!\text{C}} = (T{\,^{\circ}!\text{F}} – 32)/1.8.
Prefix relationships: 1\,\text{km}=10^{3}\,\text{m}, 1\,\mu\text{m}=10^{-6}\,\text{m}, 1\,\text{ng}=10^{-9}\,\text{g}, etc.

Problem-Solving Checklist

Clarify what is given & what is required (units!).
Sketch conversion path (factor-label strategy).
Insert conversion factors so units cancel systematically.
Compute, apply sig-fig rules, round properly.
Generate ball-park estimate; if answer is wildly different, re-check steps/units.